Difficulty: Introductory | Prerequisites: Basic algebra, familiarity with SI units (kilograms, metres, seconds)
Big picture: This section lays the thermochemistry groundwork you will lean on for the rest of organic chemistry. Before you can talk about why reactions happen, you need a shared vocabulary for energy, heat, and temperature, and you need to know the maths that connects them. If you have not touched chemistry since secondary school, start here. Everything from bond energies to reaction spontaneity (later chapters) assumes you are comfortable with these definitions and formulas.
Energy is the capacity to do work or supply heat, and it is always conserved. Heat is how thermal energy moves between objects at different temperatures. Specific heat capacity tells you how much energy it takes to warm a given mass of a substance by one degree, and the formula q = c · m · ΔT lets you calculate heat gained or lost.
Energy
The capacity to supply heat or do work. In chemistry, it is what drives reactions and determines whether products or reactants are favoured.
In simple terms, energy is whatever lets something change, move, or warm up.
Work (w)
The process of causing matter to move against an opposing force.
Think of it as: if you push something and it moves, you have done work on it.
Potential energy
The energy an object has because of its relative position, composition, or condition.
In simple terms, it is stored energy. A ball at the top of a hill has it; so does a chemical bond that has not yet broken.
Kinetic energy (KE)
The energy an object possesses because of its motion.
Think of it as: anything that is moving carries kinetic energy. Faster motion means more of it.
Law of Conservation of Energy (First Law of Thermodynamics)
During a chemical or physical change, energy can neither be created nor destroyed; it only changes form.
In simple terms, the total energy before a change equals the total energy after. Nothing is lost, nothing appears from nowhere.
Thermal energy
Kinetic energy associated with the random motions of atoms and molecules.
Think of it as the "jiggle energy" inside a substance. More jiggling means more thermal energy.
Temperature
A quantitative measure of how "hot" or "cold" something is, reflecting the average kinetic energy of its particles.
In simple terms, temperature is a number that tells you how fast, on average, the atoms in a substance are vibrating.
Heat (q)
The transfer of thermal energy between two bodies at different temperatures.
Think of it as energy in transit. Heat is not something a substance "has"; it is what flows from the hotter object to the cooler one until they match.
Exothermic process
A change that releases heat to the surroundings.
In simple terms, the system warms up its environment. A burning match is exothermic.
Endothermic process
A reaction or change that absorbs heat from the surroundings.
In simple terms, the system pulls energy in, so the surroundings cool down. An ice pack activating is endothermic.
Calorie (cal)
The amount of energy required to raise one gram of water by 1 °C (1 K). The exact value depends on atmospheric pressure and starting temperature.
In simple terms, it is an older energy unit still used in food science, but chemistry now prefers joules.
Joule (J)
The SI unit of energy. 1 J = 1 kg·m²/s², also called 1 newton-metre. Named after the English physicist James Prescott Joule.
1 kilojoule (kJ) = 1,000 J
1 cal = 4.184 J
Heat capacity (C)
The quantity of heat (q) a body of matter absorbs or releases when it experiences a temperature change (ΔT) of 1 °C or 1 K. It is an extensive property, meaning its value scales with the amount of substance present.
In simple terms, it tells you how much energy the whole object needs to warm up by one degree. A bathtub of water has a much larger heat capacity than a teacup of water.
Specific heat capacity (c)
The quantity of heat required to raise the temperature of one gram of a substance by 1 °C (1 K). It is an intensive property, meaning it does not depend on how much substance you have.
In simple terms, it is the "per-gram" version of heat capacity. Water's specific heat is 4.184 J/(g·°C), which is unusually high, and that matters a lot in climate and biology.
Energy comes in two broad types: potential energy (stored) and kinetic energy (motion).
Energy can convert freely between forms, but the total is always conserved. This is the First Law of Thermodynamics.
When one substance converts to another, energy changes form. Heat is the most common form released or absorbed, but light and electrical energy also appear in some conversions.
The Law of Conservation of Matter is a parallel idea: during a chemical change, there is no detectable change in the total amount of matter.
An object is "hot" when its atoms vibrate faster than normal (higher average KE). It is "cold" when they vibrate more slowly (lower average KE).
Adding thermal energy raises the temperature; removing it lowers the temperature.
Most substances expand when heated and contract when cooled. Thermometers exploit this: mercury/alcohol-based and bimetallic-strip thermometers both work on this principle.
Heat (q) flows from the hotter body to the cooler body and continues until both reach the same temperature (thermal equilibrium).
Heat capacity (C) is extensive: it depends on how much substance you have.
Specific heat capacity (c) is intensive: it is a per-gram property of the material itself.
If you know a substance's mass and its specific heat, measuring the temperature change before and after heating (or cooling) lets you calculate the heat transferred.
If a substance loses thermal energy, the final temperature is lower than the initial, so ΔT < 0 and q is negative.
If a substance gains thermal energy, the final temperature is higher, so ΔT > 0 and q is positive.
Heat capacity:
C = q / ΔT
Where C is heat capacity, q is heat transferred, and ΔT is the change in temperature.
Specific heat capacity:
c = q / (m · ΔT)
Where c is specific heat capacity, m is mass in grams, and ΔT is the temperature change.
Calculating heat transferred:
q = c · m · ΔT = c · m · (T_final − T_initial)
This is the workhorse formula of this section. You will use it repeatedly in calorimetry problems.
Water's high specific heat capacity is the reason coastal climates are milder than inland ones: the ocean absorbs and releases enormous amounts of heat with only modest temperature swings. The same property makes water an excellent coolant in car radiators and industrial systems.
Calorimetry (measuring heat with q = c · m · ΔT) is the basis of nutritional labels. The "Calories" on food packaging are kilocalories, each equal to 4,184 J.
Students often treat heat and temperature as the same thing. They are not. Temperature is a measure of average kinetic energy; heat is energy being transferred between objects at different temperatures.
Students frequently forget that heat capacity is extensive (depends on amount) while specific heat capacity is intensive (a material property). Mixing these up leads to wrong setups in calorimetry problems.
A negative q does not mean "no energy." It means the substance released energy. The sign tells you the direction of flow, not the magnitude.
Students sometimes assume ΔT means "final temperature." It does not. ΔT = T_final − T_initial, and the sign matters.
⚠️ The formula q = c · m · ΔT is almost guaranteed to appear on exams. Know it cold, and pay attention to sign conventions.
⚠️ Be able to convert between joules and calories (1 cal = 4.184 J). Unit conversion errors are the most common way to lose marks on heat problems.
⚠️ Know the difference between heat capacity (C, extensive) and specific heat capacity (c, intensive). Exam questions frequently test whether you can identify which one to use.
⚠️ Expect questions that ask you to determine whether a process is exothermic or endothermic based on the sign of q or ΔT.
True or false: Energy can be created during a highly exothermic reaction. False. Energy is conserved; it changes form but is never created or destroyed.
Fill in the blank: Heat flows from the ______ body to the ______ body until thermal equilibrium is reached. Hotter; cooler.
True or false: Specific heat capacity depends on how much of the substance you have. False. Specific heat capacity is an intensive property. Heat capacity (without "specific") is the extensive one.
Fill in the blank: 1 cal = ______ J. 4.184 J.
True or false: If ΔT is negative, the substance absorbed heat. False. A negative ΔT means the substance lost thermal energy (cooled down).
Q: A 50.0 g sample of a metal absorbs 250 J of heat and its temperature rises from 20.0 °C to 30.0 °C. What is the specific heat capacity of the metal?
A: c = q / (m · ΔT) = 250 J / (50.0 g × 10.0 °C) = 0.500 J/(g·°C).
Q: Is boiling water an endothermic or exothermic process? Explain using the definition.
A: Endothermic. The water absorbs heat from its surroundings (the stove) to undergo the phase change from liquid to gas.
Q: A reaction releases 5,000 J of energy. What is the sign of q, and how many kilocalories is that?
A: q is negative (energy is released, so the system loses heat). 5,000 J ÷ 4,184 J/kcal ≈ 1.20 kcal.
Q: Why is heat capacity called an extensive property?
A: Because its value depends on the amount of substance. Double the mass and you double the heat capacity. Specific heat capacity, by contrast, is the same regardless of sample size.
Q: Two objects are in thermal contact. Object A is at 80 °C and Object B is at 20 °C. Describe what happens and when it stops.
A: Heat flows from Object A (hotter) to Object B (cooler). A's temperature falls and B's temperature rises until both reach the same equilibrium temperature. At that point, net heat transfer stops.
This material connects directly to calorimetry (Chapter 5, later sections), where you will use q = c · m · ΔT in experimental setups like coffee-cup and bomb calorimeters.
The concept of exothermic and endothermic processes will underpin your understanding of enthalpy (ΔH) and Hess's Law.
Understanding the sign of q sets you up for Gibbs free energy (ΔG = ΔH − TΔS), where the direction of energy flow helps predict whether a reaction is spontaneous.
energy, work, potential energy, kinetic energy, law of conservation of energy, first law of thermodynamics, thermal energy, temperature, heat, heat transfer, exothermic, endothermic, calorie, kilocalorie, joule, kilojoule, heat capacity, specific heat capacity, q = cmΔT, calorimetry, ΔT, sign convention, extensive property, intensive property, thermochemistry, organic chemistry energy basics, Chapter 5.1